Water can act both as an oxidizing agent, in which case the water is reduced to H2, and a reducing agent, with the attendant production of O2. If a chemical species is to be stable in aqueous solution, it must not react with the water through a redox process. In other words, the oxidation and reduction potentials of the species must be such that it is thermodynamically unfavorable for the species to be either oxidized or reduced by the water.
In acid solution, the primary redox process occurs between H+ and metals producing metal ions and H2. The potential for the reduction of H+ to H2 is given by
(assuming H2 gas is at 1 atm.)
When water acts as a reducing agent, the relevant half reaction is
where E is given by
Remember that this is a reduction potential, so for water to be oxidized, the species that is going to be reduced must have a reduction potential that is more positive than 1.23V at pH=0. This requires a fairly strong oxidizing agent.
In practice it turns out that the reduction potential referred to above needs to be at least 0.6V higher than predicted due to a kinetic phenomenon known as the overvoltage (or overpotential). This "extra energy demand" is always found when gases are involved in the redox processes, and is also associated with the complexities of transferring 4 electrons at once, as in the oxidation of water.
The diagram shown here shows how the potentials for reduction and oxidation of water vary with pH. These are the inner two lines that slope downward from low pH to high pH. Note that the pH scale only runs from 2-10. For both oxidation and reduction of water, an additional line is shown that lies 0.6V above (for oxidation of water) or below (for reduction) the theoretical E. This pair of lines represents the potentials including an approximation for the overvoltage. Lastly, there is a pair of vertical lines at pH=4 and 9. These are reflective of the fact that most natural waters have a pH somewhere between these limits. The second plot emphasizes the region bounded by the upper and lower potential lines and pH 4 and 9. This region is frequently referred to as the stability field of natural waters.
A Pourbaix diagram is an attempt to overlay the redox and acid-base chemistry of an element onto the water stability diagram. The data that are required are redox potentials and solubility products. Below is the Pourbaix diagram for iron that we looked at in class) the reproduction is not very good). Below that is the same diagram showing only those species stable between the water limits, and the third diagram emphasizes the region of normal natural waters. Looking at these three diagrams can you explain why in the FeSCN2+ experiment all of the solutions were prepared in fairly concentrated nitric acid?
The next diagram shows approximate redox and acid-base behavior for various types of natural waters. You can use this diagram in conjunction with the Pourbaix diagrams above to explain and predict a number of phenomena. For example, in Lake Waban and similar bodies of water, vertical mixing of material in the lake is greatly retarded by temperature gradients. At the surface, you would expect to find Fe present mostly as Fe2O3(s), but as the solid settled to the bottom, where there is far less oxygen, and probably organic matter from decay, it is likely that the iron will be reduced to Fe2+. The ions migrate upwards and are reoxidized to Fe2O3(s).
Here are several other Pourbaix diagrams (again, the reproductions are not great). We will use these in class to answer a variety of questions. This is a large diagram so you will need to scroll in order to find the element you want.
Here are a few Pourbaix related problems:
1. Discuss the chemistry of the FeO42- ion in aqueous solution.
2. What phosphorous oxidation states are stable in water?
3. Verify by calculation the position of the line separating Fe2+ and Fe(OH)3(s).
4. List those elements that have only one form that is stable in aqueous solution.
5. List at least five elements that have quite complex aqueous solution chemistries.
|William F. Coleman firstname.lastname@example.org|
|Dept. of Chemistry, Wellesley College|
|Date Created: February 22, 2003|
|Last Modified: October 24, 2005|
|Expires: September 15, 2014|
|Copyright 2003 by William F. Coleman|